Remembering General Chemistry: Electronic Structure and Bonding

What Is Organic Chemistry?

Organic chemistry is the study of compounds primarily based on carbon. Historically, organic compounds were thought to originate only from living organisms, while inorganic compounds came from minerals. Today, organic compounds are defined as carbon-based molecules, regardless of their origin.

• Early definition: Organic compounds were believed to possess a 'vital force' from living organisms.

• Current definition: Organic compounds are any compounds containing carbon atoms.

• Example: The synthesis of urea from ammonium cyanate demonstrated that organic compounds can be made from inorganic sources.

What Makes Carbon So Special?

Carbon is unique because it can share electrons, forming stable covalent bonds. Atoms to the left of carbon in the periodic table tend to lose electrons, while those to the right tend to gain electrons. Carbon's ability to share electrons allows it to form a wide variety of compounds.

The Structure of an Atom

Atomic Structure

An atom consists of a nucleus (containing protons and neutrons) surrounded by an electron cloud. The atomic number indicates the number of protons in the nucleus, which also equals the number of electrons in a neutral atom.

• Protons: Positively charged particles in the nucleus.

• Neutrons: Neutral particles in the nucleus.

• Electrons: Negatively charged particles in the electron cloud.

• Example: Carbon has an atomic number of 6, so it has 6 protons and 6 electrons.

Isotopes

Isotopes are atoms of the same element with the same atomic number but different mass numbers due to varying numbers of neutrons.

• Example: Carbon has three common isotopes: 12C, 13C, and 14C.

The Distribution of Electrons in an Atom

Electron Shells and Orbitals

Electrons occupy shells around the nucleus, with the first shell being closest and lowest in energy. Within each shell, electrons fill atomic orbitals according to specific rules.

• Aufbau Principle: Electrons fill the lowest energy orbitals first.

• Pauli Exclusion Principle: No more than two electrons can occupy the same orbital.

• Hund’s Rule: Electrons fill empty degenerate orbitals before pairing up.

Periodic Trends: Electron Loss and Gain

Atoms in the First Column Lose Electrons

Elements like lithium and sodium, found in the first column of the periodic table, achieve stable electron configurations by losing one electron, forming positively charged ions.

• Example:

• Example:

Atoms on the Right Side Gain Electrons

Elements such as fluorine and chlorine gain electrons to fill their outer shells, forming negatively charged ions.

• Example:

• Example:

Hydrogen: Electron Loss or Gain

Hydrogen can either lose an electron to become a proton or gain an electron to become a hydride ion, achieving stability in both cases.

• Example: (proton)

• Example: (hydride ion)

Achieving a Filled Outer Shell: Covalent Bonding

Electron Sharing and Covalent Bonds

Atoms can achieve filled outer shells by sharing electrons, forming covalent bonds. This is common for nonmetals such as fluorine and hydrogen.

• Covalent bond: A bond formed by the sharing of electrons between atoms.

• Example: Two fluorine atoms share electrons to form .

• Example: Two hydrogen atoms share electrons to form .

How Many Bonds Does an Atom Form?

The number of bonds an atom forms depends on its valence electrons and the need to achieve a stable configuration. Oxygen forms two bonds, nitrogen three, carbon four, and hydrogen one.

• Example: Water (), ammonia (), methane ().

Phosphorus and Sulfur: Expanded Octet

Phosphorus and sulfur can have more than eight electrons in their outer shell due to the presence of d orbitals, allowing them to form expanded octets.

• Example: Phosphorus can be surrounded by 10 electrons; sulfur can be surrounded by 10 or 12 electrons.

Bond Polarity and Electronegativity

Nonpolar and Polar Covalent Bonds

Bonds can be classified as nonpolar or polar based on the difference in electronegativity between the bonded atoms. Nonpolar bonds occur when atoms have similar electronegativities, while polar bonds occur when there is a significant difference.

• Nonpolar covalent bond: Electrons are shared equally.

• Polar covalent bond: Electrons are shared unequally.

• Example: (nonpolar), (polar).

Electronegativity Difference and Bond Polarity

The greater the difference in electronegativity between two atoms, the more polar the bond. If the difference is very large, the bond becomes ionic.

• Electronegativity: The ability of an atom to attract electrons in a bond.

• Bond types: Nonpolar covalent, polar covalent, ionic.

Dipole Moment

The dipole moment of a bond is a measure of its polarity, depending on the difference in electronegativity and the distance between charges.

Electrostatic Potential Maps

Electrostatic potential maps visually represent the distribution of electron density in molecules, highlighting regions of positive and negative charge.

Lewis Structures and Formal Charge

Lewis Structures

Lewis structures show how atoms are bonded and indicate lone pairs and formal charges. They are essential for understanding molecular structure and reactivity.

• Lone pairs: Non-bonding pairs of electrons.

• Formal charge: The charge assigned to an atom in a molecule, calculated by comparing the number of valence electrons in the free atom to those assigned in the molecule.

Bonding Patterns of Common Elements

• Carbon: Forms 4 bonds; if not, it carries a charge.

• Nitrogen: Forms 3 bonds and has 1 lone pair; deviations result in a charge.

• Oxygen: Forms 2 bonds and has 2 lone pairs; deviations result in a charge.

• Halogens: Form 1 bond and have 3 lone pairs; hydrogen forms 1 bond.

Bond and Lone Pair Rule

The sum of the number of bonds and lone pairs for second-row elements (C, N, O, F) always equals four.

How to Draw a Lewis Structure

To draw a Lewis structure:

1. Determine the total number of valence electrons.

2. Arrange atoms and connect with single bonds.

3. Distribute remaining electrons as lone pairs.

4. Complete octets and check for formal charges.

Alternative Structural Representations

Kekulé, Condensed, and Skeletal Structures

Organic molecules can be represented in several ways:

• Kekulé structures: Show all bonds but omit lone pairs.

• Condensed structures: Omit some or all bonds, showing groups of atoms.

• Skeletal structures: Show carbon-carbon bonds as lines, omitting carbons and hydrogens bonded to carbons.

Atomic Orbitals and Molecular Bonding

s and p Atomic Orbitals

Atomic orbitals are regions around the nucleus where electrons are likely to be found. The s orbital is spherical, while p orbitals have two lobes with opposite phases.

• s orbital: Spherical shape.

• p orbital: Dumbbell shape with a nodal plane.

Bond Formation: Sigma and Pi Bonds

Bonds are formed by the overlap of atomic orbitals. Sigma (σ) bonds result from end-to-end overlap, while pi (π) bonds result from side-to-side overlap of p orbitals.

• Sigma bond: Strongest type of covalent bond.

• Pi bond: Weaker, formed in double and triple bonds.

Hybridization

Hybridization is the mixing of atomic orbitals to form new hybrid orbitals, which determine molecular geometry and bond angles.

• sp3 hybridization: Tetrahedral geometry, bond angle ≈ 109.5°.

• sp2 hybridization: Trigonal planar geometry, bond angle ≈ 120°.

• sp hybridization: Linear geometry, bond angle ≈ 180°.

Bond Strength, Length, and Order

Bond Order

Bond order refers to the number of bonds between two atoms. Higher bond order means stronger and shorter bonds.

• Single bond: Bond order 1.

• Double bond: Bond order 2.

• Triple bond: Bond order 3.

Bond Strength and Length

The more bonds holding two atoms together, the stronger and shorter the bond. Hybridization also affects bond strength and length; more s character results in shorter, stronger bonds.

• Example: Triple bonds (sp) are shorter and stronger than double (sp2) or single (sp3) bonds.

Dipole Moments and Molecular Polarity

Dipole Moments of Molecules

The dipole moment of a molecule depends on the magnitude and direction of individual bond dipoles. Symmetrical molecules do not have a net dipole moment.

• Example: Methane is symmetrical and nonpolar; water is bent and polar.

Learning Objectives

• Write ground-state electronic configurations for elements hydrogen through calcium.

• Describe relative polarity of bonds and determine dipole directions.

• Represent organic compounds using Lewis, Kekulé, condensed, and skeletal structures.

• Assign lone pairs and calculate formal charges.

• Determine hybridization of C, N, or O atoms from molecular formula.

• Describe how molecular geometry is determined by hybridization.

• Explain how hybridization affects bond strength and length.

• Describe how bond order affects bond length and strength.